Question

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in \(\mathrm{O}_{3},(\mathrm{~b})\) phosphorus in \(\mathrm{PF}_{6}^{-}\), (c) nitrogen in \(\mathrm{NO}_{2}\), (d) iodine in \(\mathrm{ICl}_{3}\), (e) chlorine in \(\mathrm{HClO}_{4}\) (hydrogen is bonded to \(\mathrm{O}\) ).

Short answer

The formal charges on the indicated atoms in the given molecules or ions are: (a) Central Oxygen atom in \(\mathrm{O}_3\): +4, (b) Phosphorus in \(\mathrm{PF}_{6}^{-}\): +1, (c) Nitrogen in \(\mathrm{NO}_{2}\): +2, (d) Iodine in \(\mathrm{ICl}_3\): +2, (e) Chlorine in \(\mathrm{HClO}_{4}\): +5.

Step-by-step solution

(a) Formal Charge on Central Oxygen Atom in \(\mathrm{O}_3\)

To calculate the formal charge on the central oxygen atom in \(\mathrm{O}_3\), we need to count the valence electrons, non-bonding electrons, and bonding electrons. Valence Electrons: Each oxygen atom has 6 valence electrons. Non-bonding Electrons: The central oxygen atom forms 2 double bonds with the other 2 oxygen atoms, so it has no electron pairs that are not involved in bonding. Bonding Electrons: The central oxygen atom forms 2 double bonds with the other 2 oxygen atoms, so it has 4 bonding electrons. Now we can calculate the formal charge using the formula: Formal Charge = Valence Electrons - Non-bonding Electrons - ½ Bonding Electrons Formal Charge = 6 - 0 - (½ × 4) Formal Charge = 6 - 2 Formal Charge = +4 So, the formal charge on the central oxygen atom in \(\mathrm{O}_3\) is +4.

(b) Formal Charge on Phosphorus in \(\mathrm{PF}_{6}^{-}\)

To calculate the formal charge on the phosphorus atom in \(\mathrm{PF}_{6}^{-}\), we need to count the valence electrons, non-bonding electrons, and bonding electrons. Valence Electrons: A phosphorus atom has 5 valence electrons. Non-bonding Electrons: Phosphorus forms 6 bonds with the fluorine atoms, so it has no electron pairs that are not involved in bonding. Bonding Electrons: Phosphorus forms 6 single bonds with the fluorine atoms, so it has 6 bonding electrons. Now we can calculate the formal charge using the formula: Formal Charge = Valence Electrons - Non-bonding Electrons - ½ Bonding Electrons Formal Charge = 5 - 0 - (½ × 6) Formal Charge = 5 - 3 Formal Charge = +2 Since the ion has an overall charge of -1, the formal charge on the phosphorus atom in \(\mathrm{PF}_{6}^{-}\) is +1.

(c) Formal Charge on Nitrogen in \(\mathrm{NO}_{2}\)

To calculate the formal charge on the nitrogen atom in \(\mathrm{NO}_{2}\), we need to count the valence electrons, non-bonding electrons, and bonding electrons. Valence Electrons: A nitrogen atom has 5 valence electrons. Non-bonding Electrons: Nitrogen forms 2 bonds with oxygen atoms, so it has 1 electron pair that is not involved in bonding. Bonding Electrons: Nitrogen forms 2 single bonds with the oxygen atoms, so it has 2 bonding electrons. Now we can calculate the formal charge using the formula: Formal Charge = Valence Electrons - Non-bonding Electrons - ½ Bonding Electrons Formal Charge = 5 - 2 - (½ × 2) Formal Charge = 5 - 2 - 1 Formal Charge = +2 So, the formal charge on the nitrogen atom in \(\mathrm{NO}_{2}\) is +2.

(d) Formal Charge on Iodine in \(\mathrm{ICl}_3\)

To calculate the formal charge on the iodine atom in ICl3, we need to count the valence electrons, non-bonding electrons, and bonding electrons. Valence Electrons: An iodine atom has 7 valence electrons. Non-bonding Electrons: Iodine forms 3 bonds with the chlorine atoms, so it has 2 electron pairs (4 electrons) that are not involved in bonding. Bonding Electrons: Iodine forms 3 single bonds with the chlorine atoms, so it has 3 bonding electrons. Now we can calculate the formal charge using the formula: Formal Charge = Valence Electrons - Non-bonding Electrons - ½ Bonding Electrons Formal Charge = 7 - 4 - (½ × 3) Formal Charge = 7 - 4 - 1.5 Formal Charge = +1.5 Since formal charges are typically whole numbers, the formal charge on the iodine atom in \(\mathrm{ICl}_3\) is most likely +2 due to a different resonance structure.

(e) Formal Charge on Chlorine in \(\mathrm{HClO}_{4}\)

To calculate the formal charge on the chlorine atom in \(\mathrm{HClO}_{4}\), we need to count the valence electrons, non-bonding electrons, and bonding electrons. Valence Electrons: A chlorine atom has 7 valence electrons. Non-bonding Electrons: Chlorine forms 4 bonds with the oxygen atoms and 1 bond with the hydrogen atom so it has no electron pairs that are not involved in bonding. Bonding Electrons: Chlorine forms 4 single bonds with the oxygen atoms and 1 single bond with the hydrogen atom, so it has 5 bonding electrons. Now we can calculate the formal charge using the formula: Formal Charge = Valence Electrons - Non-bonding Electrons - ½ Bonding Electrons Formal Charge = 7 - 0 - (½ × 5) Formal Charge = 7 - 2.5 Formal Charge = +4.5 Since formal charges are typically whole numbers, the formal charge on the chlorine atom in \(\mathrm{HClO}_{4}\) is most likely +5 due to a different resonance structure.

Key concepts

Lewis Structures

Lewis Structures serve as a visual representation of molecules depicting how atoms are connected. Each bond in a Lewis structure represents a pair of shared electrons. These structures are essential as they provide insight into the arrangement of atoms and help predict molecular geometry and reactivity. Drawing a Lewis structure involves:

• Identifying the total number of valence electrons available in the molecule.
• Arranging these electrons to satisfy the octet rule, which states that atoms tend to form bonds until they have eight electrons in their outer shell.
• Placing lone pairs or non-bonding electrons around the atoms to complete their valence shells.

By understanding how to create these structures, you deepen your understanding of molecule connectivity and predict how they might interact with other molecules.

Molecular Geometry

Molecular Geometry describes the three-dimensional arrangement of atoms within a molecule. It plays a critical role in determining physical and chemical properties, such as polarity, color, and phase of matter. The geometry is determined by:

• The number of bonding pairs around a central atom.
• Any lone pairs of electrons which may distort the shape.

Common geometrical shapes include linear, bent, trigonal planar, tetrahedral, among others. For example, ozone (O\(_3\)) has a bent shape, while phosphorus pentachloride (PF\(_6^{-}\)) has an octahedral geometry. Understanding these shapes helps predict bond angles and molecular behavior.

Valence Electrons

Valence Electrons are the electrons present in the outermost shell of an atom. These electrons are pivotal in chemical bonding since they are the electrons involved in forming bonds. Important points regarding valence electrons include:

• The number of valence electrons determines the reactivity of an element.
• They are used in the formation of covalent bonds by sharing between atoms.
• Valence electrons in ions can be different due to loss or gain of electrons in the course of forming the ion.

For example, in a nitrogen atom within NO\(_2\), there are 5 valence electrons, which account for forming its bonds with oxygen. Knowing how to determine the number of valence electrons enables you to predict how an element might form bonds in a compound.

Chemical Bonding

Chemical Bonding is the attractive force that holds atoms together in molecules. Understanding bonding is crucial as it tells why atoms combine and determines molecule stability and reactivity. The primary types of chemical bonding include:

• Covalent Bonds: Involve sharing of electrons between atoms. For example, in ozone (O\(_3\)), oxygen atoms share electrons to form covalent bonds.
• Ionic Bonds: Form through the transfer of electrons, resulting in attraction between positively and negatively charged ions.
• Metallic Bonds: Occur between metal atoms, where electrons are free to move throughout the structure.

Grasping these concepts allows you to classify compounds and appreciate properties such as melting points and solubility.